1. Liquids and solids

2. Properties of Liquids and the Kinetic molecular theory
Fluids
Substances that can flow and therefore take the shape of their container

3. Properties of Liquids and the Kinetic molecular theory
Relative High Density
10% less dense than solids (average)
Water is an exception
1000X more dense than gas
Intermolecular forces hold particles together

4. Properties of Liquids and the Kinetic molecular theory
Relative Incompressibility
The volume of liquids doesn’t change appreciably when pressure is applied
Example: Brake Fluid

5. Properties of Liquids and the Kinetic molecular theory
Ability to diffuse
Liquids diffuse and mix with other liquids
Rate of diffusion increases with temperature ( average Kinetic Energy)
Occurs because the particles are in constant motion

6. Properties of Liquids and the Kinetic molecular theory
Surface Tension
A force that tends to pull adjacent parts of a liquid’s surface together, thereby decreasing surface area to the smallest possible size
Hydrogen bonding in water creates stronger than normal surface tension
Capillary Action
The attraction of the surface of a liquid to the surface of a solid
Meniscus – the liquid is attracted to the glass and is pulled upward into the test tube creating a concave surface

7. Properties of Liquids and the Kinetic molecular theory
Evaporation and Boiling
Vaporization
The process by which a liquid or solid changes to a gas
Evaporation
The process by which particles escape from the surface of a non- boiling liquid and enter the gas state
Evaporation is a form of vaporization
Boiling
The change of a liquid to bubbles of vapor that appear throughout the liquid

8. Properties of Liquids and the Kinetic molecular theory
Formation of Solids
Freezing (or Solidification)
The physical change of a liquid to a solid by removal of heat

9. Properties of Solids and the Kinetic molecular theory
Types of Solids
Crystalline Solids
Substances in which the particles are arranged in an orderly, geometric, repeating pattern
Amorphous Solids
Substances in which the particles are arranged randomly

10. Properties of Solids and the Kinetic molecular theory
Definite Shape and Volume
Particles are packed closely together

11. Properties of Solids and the Kinetic molecular theory
Definite Melting Point
Melting is the physical change of a solid to a liquid by the addition of heat
Melting point is the temperature at which a solid becomes a liquid
Crystalline solids have definite melting points
Amorphous solids do not have definite melting points

12. Properties of Solids and the Kinetic molecular theory
High Density and Incompressibility
Due to the particles being packed closely together

13. Properties of Solids and the Kinetic molecular theory
Low Rate of Diffusion
Two solids in contact will experience VERY SLOW rates of diffusion
Particles are in a relatively fixed position and are unable to flow

14. Crystalline Solids Crystal structure
The total three dimensional arrangement of particles of a crystal
Monoclinic
Hexagonal
Orthorhombic
Trigonal
Tetragonal
Triclinic

15. fluorite

16. azurite

17. Emerald

18. Apophyllite

19. Baryte

20. Crystalline Solids Unit Cell
The smallest portion of a crystal lattice that shows the three- dimensional pattern of the entire lattice

21. Crystalline Solids Ionic crystals
Have a regular arrangement of charged particles, melt at a high temperature, are brittle, and are good insulators
Ex: Table Salt

22. Crystalline Solids Covalent network crystals
Have a large number of atoms covalently bonded to one another, melt at a high temperature, and tend to be nonconductors or semiconductors
Ex: Diamond, Quartz

23. Crystalline Solids Metallic crystals
Have metal atoms surrounded by a sea of valence electrons and are good conductors
Ex: Pyrite (fool’s gold)

24. Crystalline Solids Covalent molecular crystals
Have covalently bonded molecules held together by intermolecular forces, have low melting points, and are good insulators
Ex: Ice

25. Amorphous Solids
“Amorphous”
Greek for “without shape”

26. Amorphous solids Formation of Amorphous solids
Rapid cooling of molten materials can prevent the formation of crystals
Glass
Obsidian

27. Changes of State Equilibrium
Dynamic condition in which two opposing changes occur at equal rates in a closed system

28. Equilibrium Equilibrium and Changes of State Phase
Any part of a system that has uniform composition and properties
Condensation
The process by which gas changes to a liquid
A closed system at constant temperature will reach an equilibrium position at which the rates of evaporation and condensation will be the same

29. Possible phase changes
Change of State Process Example
Solid  Liquid Melting Candle
Solid Gas Sublimation Dry Ice
Liquid  Solid Freezing Lake Freezing
Liquid  Gas Vaporization Puddle
Gas  Liquid Condensation Coke “sweating”
Gas  Solid Deposition Frost

30. Equilibrium
An Equilibrium Equation
Liquid + heat energy vapor

31. Equilibrium Le Chatelier’s Principle
When a system at equilibrium is disturbed by application of a stress, it attains a new equilibrium position that minimizes the stress
Change in Concentration, Pressure, or Temperature
This will cause a shift in equilibrium
Liquid + Heat Energy Vapor

32. Equilibrium Equilibrium and Temperature
Increasing the temperature will move more particles into the vapor phase to compensate for the new energy

33. Equilibrium Equilibrium and Concentration
If the mass and temperature of a system remain constant, but the volume of the system increases, equilibrium will shift in order to maintain the concentrations of vapor particles

34. Equilibrium vapor pressure of a liquid
The pressure exerted by a vapor in equilibrium with its corresponding liquid at a given temperature

35. Equilibrium vapor pressure of a liquid
Volatile Liquids
Liquids that have weak forces of attraction and evaporate easily
Examples:
Alcohol
Petrol
Methylated Spirits
Acetone
Chloroform
Perfume

36. Equilibrium vapor pressure of a liquid
Nonvolatile Liquids
Liquids that have strong forces of attraction and do not evaporate easily
Examples:
Water
Mercury
Oil
Tar
Glycerine

37. Boiling
Boiling
The conversion of a liquid to a vapor within the liquid as well as at its surface. It occurs when the equilibrium vapor pressure of the liquid equals the atmospheric pressure
Occurs when equilibrium vapor pressure = atmospheric pressure

38. Boiling
Boiling Point
The temperature at which the equilibrium vapor pressure of the liquid equals the atmospheric pressure
Water boils at 100oC at 1 atm pressure
Water boils above 100oC at higher pressures
Water boils below 100oC at lower pressures
High elevations
Lower atmospheric pressure
Water boils at lower temperatures
Food takes longer to cook

39. Boiling Molar Heat of Vaporization
The amount of heat energy required to vaporize one mole of a liquid at its boiling point
Strong attractive forces between particles result in high molar heat of vaporization

40. Freezing and Melting Freezing Point
The temperature at which the solid and liquid are in equilibrium at 1 atm
For pure crystalline solids, the melting point and freezing point are the same
Temperature remains constant during a phase change
Liquid Solid + Heat Energy

41. Freezing and Melting Molar Heat of Fusion
The amount of heat energy required to melt one mole of solid at its melting point

42. Freezing and Melting Sublimation and Deposition
Sublimation is the change of state from a solid directly to a gas
Dry ice  Gaseous CO2
Deposition is the change of state from a gas directly to a solid

43. Phase diagrams
Graph of pressure vs. temperature that shows the conditions under which the phases of a substance exist. (Pressure is on a logarithmic scale)
Triple Point
Indicates the temperature and pressure conditions at which the solid, liquid, and vapor of a substance can coexist at equilibrium
Critical Point
Indicates the critical temperature and critical pressure
Critical Temperature
Temperature above which substance cannot exist as a liquid, regardless of pressure
Critical Pressure
Lowest pressure at which the substance can exist as a liquid at the critical temperature

46. Water
Structure
Two atoms of hydrogen and one atom of oxygen united by polar-covalent bonds
Bonds are bent at a 105o angle

47. Structure of water
The molecules in solid or liquid water are linked by hydrogen bonds

49. The expansion freezing water exerts sufficient force to fracture rock, and is a significant cause of rock weathering.

50. Water Physical Properties Boiling Point = 100oC Melting Point = 0oC
Density of Ice (0oC) = g/cm3
Density of Water (0oC) = g/cm3
Point of Maximum Density = 3.98oC
Molar Heat of Fusion = kJ/mole
Molar Heat of Vaporization = kJ/mole